CHEM 1312H 1st EditionExam # 4 Study Guide Lectures: 31 - 37Lecture 31 (4/1)Key concepts: solubility ICE tables, Ksp, Common Ion EffectReview how to do ICE tables and how the addition of a common ion effects the solubility product constant and the reaction!Lecture 32 (4/3)To recap the the rest of Ch. 17 on solubility:17.4 SOLUBILITY PRODUCT CONSTANTThe Solubility Product Constant, Ksp- a saturated solution is one in which the solution is in contact with un-dissolved soluteo the extent to which the dissolution reaction occurs is expressed by the magnitude of the equilibrium constant, which expresses how soluble a solid is in watero Ksp writing rules are the same for Kc rules (products over reactants, no solids, etc.)o if Ksp is small, only a small amount of the solid dissolves in water at a certain temperatureSolubility and Ksp- solubility: the quantity that dissolves to form a saturated solution (typically g/L)o solubility of a substance changes considerably in response to a number of factors; pH of solution, concentrations of other ions in a solution (especially common ions)- Ksp: the equilibrium constant for the equilibrium between an ionic solid and its saturated solution (no units)o Ksp is one set value for a given solute at any one temperature17.5 FACTORS THAT AFFECT SOLUBILITY1) Presence of a Common IonIf you have CaF2 (s)  Ca2+ (aq) + 2F- (aq)the presence of either Ca2+ or F- will shift the equilib. concentrations to the left and reduce the overall solubility of CaF22) pHThe solubility of any ionic compound that contains a basic anion (the anion of a weak acid) increases as the solution becomes more acidic (as more H+ is present). For instance, Mg(OH)2(s)) becomes a lot more soluble as the acidity of the solution increases.3) Formation of Complex IonsThese notes represent a detailed interpretation of the professor’s lecture. GradeBuddy is best used as a supplement to your own notes, not as a substitute.For example; AgCl’s solubility is increased in the presence of NH3 because NH3 is a Lewis base, which attracts the Ag+ ions that form after AgCl dissolves. The presence of NH3 drives the dissolution of AgCl to the right.In order for a Lewis base to increase the solubility of a metal salt, the base must be able to interact more strongly with the metal ion than water can. Something like Ag(NH3)2+ is a complex ion, which are very soluble in water. The equilibrium constant for the formation of a complex ion is called Kf, and its magnitude tells us the stability of the complex ion in aqueous solution.Other common Lewis bases that increase metal salt solubility are CN- and OH-.- Amphoterismo amphoteric oxides and hydroxides are soluble in strong acids and bases because they’recapable of behaving as either an acid or base themselveso examples: Al3+, Cr3+, Zn2+, Sn2+o metal hydroxides such as Ca(OH)2 and Fe(OH)3 can dissolve in acidic solution but can’t react with excess base and therefore aren’t amphoteric17.6 PRECIPITATION AND SEPARATION OF IONS- How can we predict whether a precipitate will form under various conditions?o Q is the reaction quotient that can determine the direction in which a rxn will proceed to reach equilibrium if Q = Ksp, the system is at equilibrium; the solution is saturated if Q < Ksp, the reaction will proceed to the right, towards the soluble ions; no preciipate will form if Q > Ksp, the reaction will shift left and a precipitate will form- Selective Precipitation of Ions: when ions are separated from each other based on the solubilitesof their saltso for ex, if a solution has both Ag+ and Cu2+, adding HCl to the solution will cause AgCl to precipitate will Cu2+ will remain in the solution because CuCl2 is still soluble17.7 QUALITATIVE ANALYSIS FOR METALLIC ELEMENTS- Qualitative analysis: determines the presence of absence of a particular metal ion relative to some threshold- Quantitative analysis: determines how much of a given substance is present- both proceed in three stages:o the ions are separated into broad groups based on solubility propertieso the ions in each group are separated by selectively dissolving members in the groupo the ions are identified by the conductance of specific testsLecture 33 (4/6)Key concepts: solutions, properties of solutions, entropySolution: a homogenous mixture of 2 or more pure substances in which no separation occurs; heterogenous is a slurry (sand in water, etc.)Solute: the substance that dissolvesSolvent: the medium in which the solute dissolvesEnergy that occurs: solute + solvent  solutionif the change in energy is negative, then the reaction is exothermic (products are more stable than reactanctsif the change in energy is positive, then the reaction is endothermic (reactants are more stable)NATURE FAVOURS EXOTHERMIC REACTIONSEntropy: a natural tendency of substances ot mix and spread into larger volumes when no restrained in some wayLecture 34 (4/6)Key concepts: colligative properties, factors of solubility, Henry’s Law, Raoult’s Law, molalityEx 1,Compare cyclohexane (C6H14) which is insoluble in H2O and glucose (C6H12O6), which is very soluble in water.- glucose has Hydrogen bonding- cyclohexane only interacts with H2O via LDF (in order to dissolve something, we need energy to break apart bonds and we don’t have enough because of poor energy)Henry’s Law: Cgas = kPgas, where C is the solubility of gas in solvent, k is Henry’s Law constant, and Pgas is the partial pressure above liquidColligative Properties- depend on the number of solute particles (not the kind of identity)- compared to pure solvents, solutions exhibit- vapor pressure- boiling point elevation- freezing point depression- osmotic pressureLecture 35 (4/10)Key concepts: osmotic pressure, colligative propertiesEx 1,You have four identical empty flasks. Into flasks A and B you add a small quantity of distilled water. Into flask C you add the same volume of seawater. Flask D you fill nearly completely with distilled water. You then place stoppers on the tops of all the flasks, and you leave all the flasks at room temperature except for flask B, which you place at a cooler temperature. After a shorter period of time, you measure the vapor pressure of all four flasks. Which would you predict to have the higher vapor pressure, A or B? A or C? A or D? Explain.A > B, because A has the higher temperatureA > C, because A is a pure solvent vs. C, which is a solutionA = BΔTf = (i)(Kf)(m)where ΔTf is Tf(solvent) –